Lecture 8: Thermochemistry Lecture 8 Topics Brown chapter 5 8.1: Kinetic vs. potential energy 5.1 8.2: Transferring energy as heat & work Thermal energy 8.3: System vs. surroundings Closed systems 8.4: First Law of Thermodynamics 5.2 Internal energy of chemical reactions Energy diagrams E, system & surroundings 8.5: Enthalpy 5.3 Exothermic vs. endothermic Guidelines thermochemical equations 5.4 Hess’s Law 5.6 8.6: Calorimetry Constant pressure calorimetry 8.7: Enthlapy of formation 5.7

Enthalpy of formation Enthalpy of formation: enthalpy to form 1 mole Enthalpies of formation can be used to calculate enthalpy of reaction. The general process of advancing scientific knowledge by making experimental observations and by formulating hypotheses, theories, and laws. It’s a systematic problems solving process AND it’s hands-on….. Experiments must be done, data generated, conclusions made. This method is “iterative”; it requires looping back and starting over if needed. [Why do you think they call it REsearch?] Often years, decades or more of experiments are required to prove a theory. While it’s possible to prove a hypothesis wrong, it’s actually NOT possible to absolutely prove a hypothesis correct as the outcome may have had a cause that the scientist hasn’t considered.

Enthalpy of formation (ΔHf) There are many forms (or expressions) of enthalpy including Hvaporization, Hfusion, Hcombination & Hformation. The most useful is standard enthalpy of formation (H°f): the enthalpy change that accompanies the formation of 1 mole of a compound from elements in their standard states. The form of the element at 1 atm, 298 K/25°C gases - 1 atm solutions - 1 M standard states? Na? Hg? N2? solid liquid gas The enthalpy of formation of any substance in its most stable form is always zero: Carbon as graphite; oxygen as O2 gas; copper as an elemental solid (metal). Write the reactions for H°f for: CH3CH2OH FePO4 2C + 1/2O2 + 3H2  CH3CH2OH Fe + P + 2O2  FePO4 see table 5.3, p. 177 & appendix C p. 183-5

Example: ΔHf For which reaction (at 25°C) would enthalpy change represent standard enthalpy of formation? If not, how could you change the reaction conditions? a. 2Na(s) + 1/2O2(g)  Na2O(s) 2K(l) + Cl2(g)  2KCl(s) C6H12O6(s)  6C(diamond) + 6H2(g) + 3O2(g) Yes - all in standard states & 1 mole product is formed No - K should be (s) & 2 moles product are formed No - this is a decomposition, not a formation; reverse it p. 183-5

ΔHfs can be summed to calculate ΔHrxn So this concept combines Hess’s Law with Hf . If we know the Hf of all reactants & products, we can calculate Hrxn: H°rxn = nH°f(products) - nH°f(reactants) Propane (C3H8) is combusted to form CO2 and H2O under standard cond. C3H8 + 5O2  3CO2 + 4H2O (gas) (gas) (gas) (liquid) H°rxn = -Hf [C3H8(g)] + 3Hf [CO2(g)] + 4Hf [H2O(l)] = -(-103.85 kJ) + 3(-393.5 kJ) + 4(-285.8 kJ) = -2220 kJ (moles)(kJ/mole) Expressed via Hess’s Law: C3H8 --> 3C + 4H2 H1 = -Hf [C3H8(g)] reversed 3C + 3O2 --> 3CO2 H2 = 3Hf [CO2(g)] stoichiometry 4H2 + 2O2 --> 4H2O H3 = 4Hf [H2O(l)] state ! C3H8 + 5O2 --> 3CO2 + 4H2O H°rxn = H1 + H2 + H3 p. 185-7

Examples: ΔHf Calculate the H for combustion of 1 mole benzene (C6H6). C6H6 + 15/2O2  6CO2 + 3H2O (liquid) (gas) (gas) (liquid) Hrxn = [6Hf(CO2) + 3Hf(H2O)] - [Hf(C6H6) + 15/2 Hf(O2)]  products  reactants = [6(-393.5 kJ) + 3(-285.8 kJ)] - [(49.0 kJ) + 15/2(0 kJ)] = (-2361 - 857.4 - 49.0) kJ = - 3267 kJ Use enthalpies of formation to calculate the Hrxn for: CaCO3(s) --> CaO(s) + CO2(g) Hrxn = [Hf(CaO) + Hf(CO2)] - [Hf(CaCO3)]  products  reactants = [(-635.5 kJ) + (-393.5 kJ)] - [(-1207.1 kJ)] = 178.1 kJ p. 185-7

Enthalpies of combustion of fuels Carbohydrates? Glucose = C6H12O6 C6H12O6 + 6O2  6CO2 + 6H2O Hrxn = - 2803 kJ Average fuel value for carbohydrates (& proteins) is ~17 kJ/g. Fats? Triacylglycerol (body fat) = C57H110O6 2C57H110O6 + 163O2 --> 114CO2 + 110H2O Hrxn = - 75,520 kJ Average fuel value for carbohydrates (& proteins) is ~38 kJ/g. What about industrial fuels? fuel value (kJ/g) Coal 31 - 32 Oil 45 Natural gas 49 Gasoline 48 H2 142 p. 188-9